ppt

```ENERGY
Energy


Energy (E) is the ability to do work.
Many types, but we can say 3 main
types:



Potential
Kinetic
US Energy Consumption by Source

Light Energy
 Visible

and Invisible
Travels in waves over
distances
 Electromagnetic
 Waves
waves
all directions from the
source
 Visible light, UV light, Infra
Potential Energy (PE)

Stored Energy
 Due
to position
Gravitational
Elastic
PE
PE
 Chemical
bonds
Chemical
PE
 Nuclear
energy
 Fuels
 Attractions
molecules
between
Kinetic Energy (KE)

Energy of motion
 Atomic
vibrations
 Molecular
movement
 Vibration
 Rotation
 Translation
 Movement
subatomic
particles
of
Kinetic Energy
Can be calculated:
How are each type shown here?

 Rainbow

Kinetic
 Windmill

= visible light
moving
Potential
 All
molecules store
energy
 Water
in clouds
 Air
 Materials
the windmill is
at the bottom
Temperature Scales: Measuring that
Thermal Energy
Boiling
Freezing
Fahrenheit (oF)
212
32
Celsius (oC)
100
0
Kelvin
373
273
A Note on the Fahrenheit Scale





NEVER use it in this class. Ever. Only Belize and the US use this
scale.
Gabriel Fahrenheit made great thermometers. His scale was
replicated the world over because of this. But if you stop and
boiling? 180 degrees separates them.
100 degrees, as in the Celcius scale (sometimes called the
Centigrade scale) makes much more sense.
Fahrenheit based 0°F on the freezing point of water mixed with
temperature (he was off by 2.6°). Why? Because he felt like it
and it was easy to draw lines at those intervals.
(According to a letter Fahrenheit wrote to his friend Herman Boerhaave,
[8] his scale was built on the work of Ole Rømer, whom he had met earlier. In Rømer’s scale, brine freezes at 0 degrees, ice
melts at 7.5 degrees, body temperature is 22.5, and water boils at 60 degrees. Fahrenheit multiplied each value by four in order to eliminate fractions and increase the granularity of the scale. He
then re-calibrated his scale using the melting point of ice and normal human body temperature (which were at 30 and 90 degrees); he adjusted the scale so that the melting point of ice would be 32
degrees and body temperature 96 degrees, so that 64 intervals would separate the two, allowing him to mark degree lines on his instruments by simply bisecting the interval six times (since 64 is 2 to
the sixth power). I took this from Wikipedia.
Kelvin Temperatures

Based on absolute zero (0 K, -273 oC)
 The
temperature at which ALL KE stops
 NO molecular motion.
 Lowest temperature theoretically possible
 Can’t
really get there in real life
3rd Law of Thermodynamics in a few
slides)
 (See

K = oC + 273
 Technically
273.14, but we can stop at 3
significant digits
Why do we need the Kelvin scale?

Two reasons

We need a scale that is relative
to molecular motion
for certain topics
 You
can’t use negative numbers to indicate motion when it
IS present
 -20°C makes NO sense in light of indicating motion
 And 40°C ISN’T twice as much motion as 20°C,


(40K IS twice the motion of 20K)
Because when working with equations, can’t use zero
We get undefined answers if we divide
 We get answers of 0 if we multiple
 And those answers would NOT make sense if compared to answers
calculated with a positive or negative number

The 4 Es:
Energy, Exergy, Entropy, & Enthalpy


Energy (E): The
ability to do work
Entropy (S): The
measure of the
disorder of a
system
There will be more on these!
Exergy: The energy
available to do
work
 No symbol


Enthalpy(H): The
thermal energy
(heat) content of a
system
Thermodynamics

The study of energy flow
 inter-relation
between heat, work, and energy of a
system

Summary of the three laws:
1.
2.
3.
The energy in the universe is constant
Things get more disorganized over time in a system
until everything is equal
You can’t reach absolute zero
st
1

Law of Thermodynamics
The energy in the universe is constant
 E=mc2
 Law

of Conservation Matter
Matter can not be created or destroyed
 Law
of Conservation of Energy
Energy
can not be created or destroyed
 However,
matter and energy can both change
forms in chemical reactions
 Can also interconvert between matter and
energy in NUCLEAR reactions (more on this later
this year.)
Summed up: You can not win. You can’t get something for nothing because
energy and matter are conserved.
Time
Energy
Before the

nd
2
Law…
Entropy (S) is a measure of
DISORGANIZATION in a system (this simply
put; there is a much more complicated
description about the unavailable energy to do
work)
 Anything
disorganized has higher entropy than
something organized

Exergy is the Energy available to do work
nd
2

Law of Thermodynamics
Things get more disorganized over time in a
system until everything equilibrium is reached
(everything is equal)
 Heat
flows from hot to cold, not the reverse
 Law of Entropy
By nature, things get more disorganized to

The quality of the energy (which is exergy)
decreases over time
Summed up: You can not break even. You can not return to the same
energy state because things get more disorganized (gain entropy)
Exergy and Energy



The energy of the universe is constant, but exergy is constantly
consumed. This can be compared with a tooth-paste tube: When you
squeeze the tube (= conduct any process) the paste (= exergy)
comes out. You can never put the paste back in the tube again (try!),
and in the end you have only the tube itself (= low-exergy) left.
When you squeeze the tube, the depressions (= entropy) will
increase. (The entropy of a system increases when exergy is lost) But
you can never take the depressions in the tube and 'un-brush' your
teeth. (I.e. entropy is not negative exergy.)
exergy. You can find the energy as room temperature heat after
some time, but you can not take that room temperature energy back
to the electricity company and ask for money back. They won't
accept it.
Energy and Matter Gain Entropy Over Time
Exergy: The Energy available to do
work
rd
3

Law of Thermodynamics
You can’t reach absolute zero and
expect things to happen
At
absolute zero, all kinetic motion
ceases. And that energy needs to go
somewhere. It goes to something else.
And gets transferred back until
everything is at an equal temperature.
Summed up: You can not get out of the game, because
absolute zero is unobtainable.
Law of Conservation of Energy

Energy cannot be created or
destroyed…but it CAN change forms.
 Example:

Burning wood in a fire
The energy in chemical bonds is released
as heat (KE and PE), light (RE), sound (KE)
 These
 have
forms of energy are less useful
less exergy
EM Waves
Potential
Energy:
Stored
Kinetic
Energy:
Motion
The CPE in these items could:
Rio Summer Olympics
Proposed Solar
Waterfall
http://www.snopes.com/ph
otos/architecture/solartowe
r.asp
Combinations of PE and KE are very
common on a large scale
KE and PE animation
PE and KE
When E changes forms…

The amount of energy one thing loses is
gained somewhere else.
E
lost = E gained (Law of Conservation of
Energy)
 But the E gained is usually not all in one
place (2nd Law of thermodynamics)
 It
 Often
in the forms of heat and light
 Which are less useful (less exergy)
Energy Transformations
Thermal Energy: KE + PE on the small scale
• What’s up with Temperature vs Heat?
• Temperature is related to the average kinetic energy
of the particles in a substance.
Thermal energy relationships
As temperature increases, so does thermal
energy (because the energy of the
particles increased).
If the temperature stays the same, the
thermal energy in a more massive
substance is higher (because it is a total
measure of energy).
Heat
Cup gets cooler while hand
gets warmer
The flow of thermal
energy from one
object to another.
Heat always
flows from
warmer to
cooler objects.
Ice gets warmer
while hand gets
cooler
Heat and Temperature
Heat: the measure of the flow of
RANDOM kinetic energy
 Temperature: the measure of heat
 So…temperature is a measure of
kinetic energy of the particles of a
substance

Thermal Energy
•Thermal Energy is the
total of all the (kinetic and
potential) heat energy of
all the particles in a
substance.
•PE from how the molecules are placed relative to
each other (attractions)
•Farther = more PE, just like how something farther
off the ground has higher gravitational PE
Exothermic and Endothermic Processes
Endothermic


Energy is being
gained/ absorbed by
the object or substance
(called the system)
from the surroundings
Have positive change
in enthalpy values
(+ΔH)
Exothermic


Energy is lost/
released from the
object or substance
(called the system) to
the surroundings
Have negative change
in enthalpy values
(-ΔH)
The big picture…

How do we see this energy cycling in the
real world, and not just as a part of
Chemistry class?
 Around
the house?
 In the environment?
 While thinking about a car?
•If the cup is the system, it is
undergoing an exothermic
process because it is losing
heat to the surroundings
(hand)
•If the ice is the system, it
is undergoing an
endothermic process
because it is absorbing
heat from the surroundings
(hand)
Cup gets cooler while
hand gets warmer
Ice gets warmer
while hand gets
cooler
Which is process is endothermic? Which
is exothermic?
Trophic Levels and Energy
Consumers
are all
heterotrophs
3°Consumers:
Carnivores and
Omnivores
2°Consumers:
Carnivores and
Omnimores
1°Consumers: Herbivores
Producers: Autotrophs
Energy Out;
90% per level
Can the world really run out of
Energy?
World-Wide Energy Sources, (2007)
PHASE
CHANGES &
ENERGY
Phase Diagrams


Tell what state of matter a material is in at a given
temperature and pressure
The triple point is the pressure and temperature when a
solid, liquid, and a gas of the same substance exist at
equilibrium

Equilibrium: When there is no net change
Here referring to changes in state
 Can also refer to temperature and chemicals



The critical point is the temperature above which a
substance will always be a gas, regardless of pressure
Phase Diagrams
Phase Diagram for Water
A few terms

Freezing Point - The temperature at which the solid
and liquid phases of a substance are in equilibrium
at atmospheric pressure.
 The


same temperature as the melting point
Boiling Point - The temperature at which the vapor
pressure of a liquid is equal to the pressure on the
liquid.
Vapor Pressure- The pressure at which the
vaporization rates are equal to condensation rates
Phase Changes
Enthalpy(H):
The heat
(thermal
energy) content
of a system
States of Matter and Entropy
The states are NOT plateaus because entropy is NOT constant.
This isn’t a phase change diagram.
Energy and Matter and Connected

Any change in
matter
ALWAYS is
accompanied
by a change in
energy
Phase Changes and Energy
Temperature, ̊C
Heating Curve
Time,
min
Why does temperature remains
constant when melting or boiling?

During melting or boiling, energy is
from the surroundings

absorbed
Due to the increase in the thermal energy of the particles
from the increase in PE of the particles
 Molecules are
moving apart
 breaking attractions which
 Absorbs latent (hidden) heat
 can not be measured on a thermometer

 Substance
(system) gets warmer
The E’s and Heating
•Endothermic process
•Energy is absorbed from surroundings
•Entropy increases
•Enthalpy is positive (+ΔH) since heat added
•Exergy decreases
Why does temperature remains constant
when freezing or condensing?

During freezing or condensing, energy is
released to the surroundings
 Due
to the decrease in the thermal energy from the
decrease in PE of the particles
 Molecules are
moving closer
 forming new attractions that are
 Releasing latent (hidden) heat


can not be measured on a thermometer
 Substance
(system) gets colder
The E’s and Cooling
•Exothermic process
•Energy is lost to surroundings
•Entropy decreases
•Enthalpy is negative (-ΔH) since heat is lost
•Exergy increases
What happens during each segment
Cooling Curve: The Reverse of a
Heating Curve
Temperature, ̊ C
Time, minutes
Measuring the Energy of Phase
Changes

The math of thermal energy flow
REMEMBER: Energy and Matter and
Connected
Any change in
matter ALWAYS
is accompanied
by a change in
energy
 This includes
changes in
temperature
and/ or phase

Specific Heat : c
• Things heat up or cool down at different
rates.
Land heats up and cools down faster than water,
and aren’t we lucky for that!?
•Specific heat is the amount of heat required to
raise the temperature of 1 kg of a material by one
degree °C
•cwater = 4.184 J / g °C
•the number is high; water “holds” its heat
•c sand= 0.664 J / g °C
•less E than water to change it; it doesn’t hold
heat as well as water does
This is why land heats up quickly during the day and
cools quickly at night and why water takes longer.
Why does water have such a high
specific heat?
water
metal
Water molecules form strong attractions with other water molecules;
it takes more heat energy to break those attractions than other
materials with weaker forces of attraction between them.
Specific Heat Capacities of Selected
Substances
 cwater
= 4.184 J / g °C
 cice = 2.09 J / g °C
 csteam = 1.99 J / g °C
 csand = 0.664 J / g °C
 cAl = 0.90 J / g °C
 cFe = 0.449 J / g °C
Heat can be Transferred even if there is No
Change in State
q = mc∆T
Remember this? Which is process is
endothermic? Which is exothermic?
Now we care about how much energy
is being transferred, and are ready to
calculate that change.
Calculating Changes in Energy: The
Calorimetry Equation
q = mcT
•q = change in thermal energy
•(+) value means heat is absorbed
•(-) value means heat is released
•m = mass of substance
•T = change in temperature (Tfinal – Tinitial)
•c = specific heat of substance
•Each substance has a different c (see CRH, p__)
•Different states of matter for the same substance may have
a different c
Specific Heat Capacity Problems
If 25.0 g of Al cool from 310 oC to 37 oC, how many joules of heat energy are lost
by the Al?
Notice that the negative sign on q signals heat “lost by” or
transferred OUT of Al.
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy changes at constant T
Ice + 334 J/g (heat of fusion) -----> Liquid water
Is there an equation? Of course!
Or… Heat Transfer can cause a
Change of State
Changes of state involve energy at constant T
H20(s) +334 J/g  H20(l)
Ice + 334 J/g (heat of fusion)  Liquid water
q = mΔHfusion
•m= mass
•ΔHfusion = the enthalpy of melting
• the change in thermal energy associated
with melting
•Units are J/g or KJ/Kg
q = mΔHfusion

WHY DO I NEED THIS WHEN I HAVE
q = mc∆T?

Well, when a phase changes THERE IS NO
change in temperature… but there is definitely
a change in energy!
Sample Problem:
How much heat energy is required to melt 25.0g of ice,
(assuming constant temperature of O°C)?
Value is positive, which means heat is absorbed, which makes
sense!
Latent heat* and the PE of particles
molecule
strong attraction
Regular
arrangement
breaks up
weak attraction
*Latent means hidden. Latent heat is the thermal energy
(potential energy) associated with the attractions between
molecules, and can not be measured with a thermometer.
Latent heat and the PE or particles
Energy has to be supplied to oppose the
attractive force of the particles.
PE  as
molecules
separate
PE related to the forces of attraction between the particles
solid  liquid or liquid  gas
average potential energy 
Latent heat and PE
The transfer of energy does not change the KE.
Temperature does not change.
latent heat = change in PE between
molecules during change of state
Video and song:
Remember……
•
Energy changes accompany changes in state; either:
•
• Gain thermal energy
• Molecules
• Move more (gain KE)
• Separate (gain PE from broken attractions between
molecules)
• Have a higher entropy
• Are more disorganized
Or
•
Energy is removed (exothermic)
• Molecules move less
• Lose thermal energy
• Move less (lose KE)
• Move closer (lose PE from new attractions between
molecules)
• Have lower entropy
• Get more organized
Latent Heats

You have a certain energy change associated with
changing state. These values are usually reported
for fusion and vaporization as:
 ΔHfusion= (latent)
Heat of fusion (melting)
 Δ Hvaporization = (latent) Heat of vaporization
 Δ Hsublimation =(latent) Heat of sublimation

Different materials have different values for each
condensation?

typically listed, but are the negative values of those
for fusion and vaporization because the energy
transferred is the same, but in the opposite direction
 (latent)
Heat of freezing= -ΔHfusion
 (latent) Heat of condensation= -Δ Hvaporization
Enthalpy changes with phase changes
Enthalpy values for H2O
∆Hfusion= 334 J/g
 ∆Hvaporization= 2259 J/g
 ∆Hsublimation = 2594 kJ/g

From:
http:/
/hype
rphysi
cs.phy
astr.gs
u.edu
/hbas
e/tabl
es/ph
ase.ht
ml#c1

http://hy
perphysic
s.phyastr.gsu.e
du/hbase
/tables/p
hase.html
#c2
Summing it all up: How do you know what to do to
calculate energy changes?
•
Check to see if there is a temperature change.
•
•
If yes, use q=mcΔT.
Also, check to see if there is a phase change.
•
If yes, you need to use
•
•
•
•
q= Δ Hfusionmass
• or
q= Δ Hvaporizationmass
depending on which one applies*
or both if there are two phase changes
*If the material freezes or condenses. You can use the negative
value Δ Hfusion or Δ Hvaporization
How much energy is required to change 0.5 kg of
water at 0 °C to ice?
Things you know:
•m =
•There is____ temperature change, and there is
change of state (freezing)
•The water is going __________
So…..
this all tells you to use _________(negative of
melting value) in
q=
q=
(The negative value makes sense since you are
cooling the water, so energy leaves)
How much energy is required to melt 0.5 kg of ice at 0 °C
temperature raised to 80 °C?
Total energy required
Heat & Changes of State
What quantity of heat is required to melt 500. g of ice and
heat the water to steam at 100. oC?
Heat of fusion of ice = 334 J/g
Specific heat of water = 4.184 J/g•°C
Heat of vaporization = 2259 J/g
+2257
J/g
+334 J/g
Putting it all together…
So… if I want the total heat to take ice and turn it to steam I need to add values from
3 steps…
1. To melt the ice I need to multiply the heat of fusion with the mass
2.
• q = ∆Hfusionm
Then, there is moving the temperature from 0°C to 100°C.
• For this there is a change in temperature so we use
•
q= mc∆T
3.
That just takes us to 100°C, what about vaporizing the molecules?
•
We need q=∆Hvaporizationm
Add up all the values, and you have it.
(However, if you are taking it from below the freezing point to above 100°C, you
need to add in the changes with q=mc ∆ T there, too!)
And now… More! Heat & Changes of State
How much heat is required to melt 500. g of ice and heat the water to
steam at 100 oC?
1.
To melt ice
2. To raise water from 0 oC to 100 oC :
3. To evaporate water at 100 oC:
4. Total heat energy =
Maybe a picture can help….
Putting it all together:

How are matter and energy related?
What influences does energy have on
matter? What does this tell us about the
world as we know it?
Making Pizza: Changing Matter
Describe the pizza making process in terms of:
 Matter
 States
(s, l, g)
 Elements, compounds, mixtures
 Homogeneous

and heterogeneous mixtures
Properties and changes
 Both
 chemical
and physical
 Intrinsic (intensive) and extrinsic (extensive)

Energy
```